Carbonation of metal silicates for long-term co2 sequestration

ABSTRACT

In a preferred embodiment, the invention relates to a process of sequestering carbon dioxide. The process comprises the steps of: (a) reacting a metal silicate with a caustic alkali-metal hydroxide to produce a hydroxide of the metal formerly contained in the silicate; (b) reacting carbon dioxide with at least one of a caustic alkali-metal hydroxide and an alkali-metal silicate to produce at least one of an alkali-metal carbonate and an alkali-metal bicarbonate; and (c) reacting the metal hydroxide product of step (a) with at least one of the alkali-metal carbonate and the alkali-metal bicarbonate produced in step (b) to produce a carbonate of the metal formerly contained in the metal silicate of step (a).

CROSS-REFERENCE TO RELATED APPLICATION

This application is a continuation of U.S. Ser. No. 12/016,098 filedJan. 17, 2008, which is a division of U.S. patent application Ser. No.10/706,583 filed Nov. 12, 2003, which claims the benefit of U.S.provisional application Ser. No. 60/464,728, filed Apr. 23, 2003.

STATEMENT REGARDING FEDERALLY SPONSORED RESEARCH OR DEVELOPMENT

The United States Government has rights in this invention pursuant toContract No. DE-AC05-00OR22725 between the United States Department ofEnergy and UT-Battelle, LLC.

BACKGROUND OF THE INVENTION

Rising levels of carbon dioxide (CO₂) in the Earth's atmosphere, causedprimarily by combustion of fossil fuels, have prompted concern thattemperatures at the Earth's surface will increase sharply during the21^(st) century. To address this issue, numerous nations are developingplans for lowering CO₂ emissions to the atmosphere. The principalapproaches under consideration are: improving energy efficiency; makinggreater use of alternative sources of energy; and developingeconomically viable technologies for capture, separation, and long-termstorage of CO₂. The latter strategy, known as “CO₂ sequestration,” isreceiving increasing attention because it permits continued use ofhigh-carbon fossil fuels to generate electrical power while ensuringthat CO₂ releases to the atmosphere are reduced.

A potentially attractive means for CO₂ sequestration is injection ofgaseous CO₂ into underground reservoirs, e.g., active or depleted oiland gas fields, deep brine formations, and subterranean coalbeds. Theunderlying premise of this approach is that, after injection, the CO₂will remain sequestered in the host rock for hundreds, perhaps eventhousands, of years. In practice, however, such long-term reservoirintegrity cannot be guaranteed. If either gaseous CO₂ or CO₂-saturatedformation water escapes to overlying strata or to the surface,underground and subaerial water supplies could become seriouslycontaminated, and/or large amounts of CO₂ could be released to theatmosphere.

Significantly, the reservoir-integrity problems associated withsubterranean sequestration of gaseous or liquid CO₂ can be completelyavoided by chemically binding CO₂ with suitable solid materials. Thisalternative CO₂ sequestration strategy, known as “mineral carbonation,”involves reaction of CO₂ with naturally occurring silicates to producesolid carbonate compounds, such as calcite (CaCO₃) and magnesite(MgCO₃), for the purpose of long-term terrestrial isolation of CO₂.“Mineral carbonation” also implies a chemical process carried out atelevated temperatures and pressures in an industrial-scale reactor,because a similar term, “mineral trapping,” alludes to crystallizationof carbonate compounds at ambient temperature and pressure after CO₂ isinjected into a subsurface geologic formation. The U.S. Department ofEnergy (DOE) classifies mineral carbonation as a “CO₂ conversion”technology, rather than a geological CO₂-sequestration strategy, becausein mineral carbonation most, if not all, of the CO₂ is converted to oneor more solid carbonate compounds, whereas in mineral trapping only atiny fraction (generally less than one volume %) of the injected CO₂ isultimately incorporated into solid carbonates.

Mineral carbonation has many important advantages over alternativemethods for large-scale CO₂ sequestration. First, the carbonatecompounds formed in the process are thermodynamically stable,environmentally benign, and weakly soluble in meteoric water.Consequently, they can be amended to soils to reduce acidity andincrease moisture content, combined with stone to strengthen roadbeds,or simply dumped in a landfill. Alternatively, the carbonates could bereturned to the site of excavation to fill the cavity created bysoil/rock removal. Regardless of the particular end use or disposalscheme selected for the carbonates, the reacted CO₂ will remain tightlybound in the crystallographic structures of the carbonates, immobilizedfor an indefinite period of time. Therefore, a commercial mineralcarbonation technology creates no major “legacy issues” for nearbypopulation centers. In contrast, other proposed methods for wide-scaleCO₂ sequestration, such as subsurface storage in brine formations, anddisposal in deep-ocean realms, rely on risky environmental factors toensure long-term CO₂ containment: an impervious, superjacent “caprock”in the case of subsurface injection of CO₂ into brine formations, andlow ambient temperature and high ambient pressure, with nocurrent-driven dispersal of the sequestration “agent” (liquid CO₂ orCO₂-hydrate), in the case of suboceanic CO₂ disposal.

In weighing the technical feasibility of CO₂ sequestration by mineralcarbonation, it should be noted that huge masses of rocks and clay-richformations suitable for carbonation occur worldwide. For example,ultramafic complexes and large serpentinite bodies are major sources ofthe magnesium-rich minerals olivine (forsterite) and serpentine, whichcan be carbonated by the reactions

Moreover, contact-metamorphosed limestones frequently containwollastonite (CaSiO₃), and large quantities of plagioclase [(Ca_(x),Na_(1−x))(Al_(1+x)Si_(3−x))O₈] are present in many different types ofcommon rocks. Wollastonite and plagioclase can be converted to calcite(plus silicious solid material) by the reactions

Another key attribute of mineral carbonation, in relation to othertechnologies that deal with CO₂ waste streams, is that costs associatedwith CO₂ transport are potentially very low. This is so because in anindustrial-scale implementation of a mineral carbonation technology, themetal-silicate feedstock can be carbonated in commercial facilitieslocated adjacent to, or near, large “point sources” of CO₂ generation,such as fossil fuel-fired power plants, cement factories, and steelmills. In contrast, CO₂ sequestration in deep brine aquifers, or thebenthic regions of the world's major oceans, would often require CO₂transport over substantial distances. Building and maintaining manymiles of pipeline to achieve such transport, or hauling liquid CO₂ overlong distances by truck, train or ship, would be extremely expensive andperhaps totally impractical.

Finally, the following additional advantages of mineral carbonation arenoteworthy: (1) by technical necessity, mineral carbonation involvesrapid conversion of CO₂ to solid carbonate(s), and (2) by virtue ofcreating one or more solid carbonate phases from a volatile phase richin CO₂, carbonate crystallization automatically produces a largereduction in total volume. It has already been demonstrated byresearchers at the Albany Research Center in Oregon, and the Los AlamosNational Laboratory in New Mexico, that, with vigorous mechanicalstirring, olivine and heat-pretreated serpentine can be quantitativelyconverted to magnesite (see Reactions 1 and 2 above) in ˜30 minutes at155° C. and 185 atm total (fluid) pressure. Significantly, the latterprocessing conditions are readily attained in modem industrial reactors.A large reduction in the total volume of the reactants (CO₂, plus one ormore condensed phases, and often one or more “additives” and/orcatalysts) is automatically achieved in mineral carbonation because theCO₂-bearing solids produced are >1000× more dense than gaseous CO₂ atSTP (standard temperature and pressure: 25° C., 1 atm). This contractionessentially eliminates the “room problem” associated with storing largevolumes of CO₂ (as a gas, liquid or supercritical fluid) in subsurfacerock formations.

While it is evident that mineral carbonation offers many importantadvantages over competing CO₂ sequestration technologies, it is alsotrue that it suffers two major disadvantages. Chief among these is theneed to mine, or quarry, large quantities of silicate feedstock tosequester the gigatons of atmospheric CO₂ generated annually bycombustion of fossil fuels. Excavating massive amounts of rock and soilto permit silicate carbonation at sites near major industrial sources ofCO₂ will be expensive, and will require intense reclamation activitiesto restore the land to an environmentally acceptable state. However,there is no doubt that this can be accomplished using modern methods ofenvironmental restoration. In addition, it is likely that newtechnologies will soon be developed to enable innovative synergies, andmore satisfactory compromises, between large-scale energy production andtraditional modes of land use.

The second major disadvantage of mineral carbonation is that elevatedtemperatures and pressures, and chemical “additives” and/or catalysts,are usually required to accelerate CO₂ conversion to one or morecrystalline carbonates. While considerable success has already beenachieved in carbonating olivine (Reaction 1) at commercially feasibletemperatures and pressures, mineral carbonation experiments performedover the past four years at the Albany Research Center have shown thatuntreated serpentine does not react as readily (Reaction 2). To date,the only known remedy for sluggish serpentine carbonation is toheat-pretreat the mineral to 600-650° C. prior to carbonation, whichdrives off structurally bound water (hydroxyl groups). Tests of thisaltered (dehydroxylated) serpentine have shown that it is much morereactive than untreated (hydroxylated) serpentine. However, at a typicalfossil fuel-fired power plant, heat treating serpentine at 600-650° C.prior to carbonation would require ˜200 kW-hr of electricity per ton ofserpentine feedstock. With one ton of carbon in a fossil fuel producing˜3.7 tons of CO₂, and each ton of CO₂ consuming ˜2.0 tons of serpentineduring carbonation, the power requirements for serpentinedehydroxylation represent 20-30% of total power output. This largeenergy penalty threatens the economic viability of CO₂ sequestration byserpentine carbonation.

It is evident from Reactions 3 and 4 that the problems plaguingserpentine carbonation would be partly or entirely avoided if a moreabundant silicate mineral could be utilized. In this regard, it isnoteworthy that wollastonite is carbonated by Reaction 3 at 60° C. usingan aqueous solution of acetic acid as a catalyst. This result is of somescientific interest, but it fails to significantly bolstermetal-silicate carbonation as a potential means for sequestering largemasses of CO₂ because wollastonite, while not rare in nature, istypically found in significant quantities only in contact metamorphicaureoles where it tends, along with other silicates, to form small,isolated bodies adjacent to igneous intrusions. The other principaloccurrence of wollastonite is as a widely disseminated mineral inregionally metamorphosed carbonate strata. Thus, wollastonite is notavailable in sufficient quantities to sustain a wide-scale silicatecarbonation technology.

The low abundance of wollastonite leaves plagioclase as the majorpotential source of calcium (Reaction 4) to produce the quantities ofcarbonate required to sequester gigatons of CO₂ by metal-silicatecarbonation. (Other, locally significant potential sources of calciuminclude Ca-rich clay deposits, Ca-rich fly ash, and waste concrete.)However, a commercially feasible plagioclase carbonation technologyfaces two formidable technical challenges. First, it is inherentlydifficult to extract calcium from plagioclase because, being a frameworksilicate with a three-dimensional structure held together by tightlybonded atoms of silicon and aluminum, plagioclase is not readilydestabilized by firing at high temperatures, or easily “digested”(decomposed) by most customary solvents. Second, while most plagioclasescontain a significant amount of calcium, Ca-contents are always lessthan that of wollastonite. Therefore, per ton of silicate feedstock,less calcium-rich carbonate (calcite) is formed from plagioclase thanfrom wollastonite. These difficulties notwithstanding, it is clear thatplagioclase carbonation merits serious scientific study to determinewhether it could be an attractive alternative to serpentine carbonationin sequestering large quantities of CO₂.

SUMMARY OF THE INVENTION

In a preferred embodiment, the invention relates to a process ofsequestering carbon dioxide. The process comprises the steps of (a)reacting a metal silicate with a caustic alkali-metal hydroxide toproduce a hydroxide of the metal formerly contained in the silicate; (b)reacting carbon dioxide with at least one of a caustic alkali-metalhydroxide and an alkali-metal silicate to produce at least one of analkali-metal carbonate and an alkali-metal bicarbonate; and (c) reactingthe metal hydroxide product of step (a) with at least one of thealkali-metal carbonate and the alkali-metal bicarbonate produced in step(h) to produce a carbonate of the metal formerly contained in the metalsilicate of step (a).

The invention also relates to a system for sequestering carbon dioxidefrom a gas stream. The system comprises a gas stream containing carbondioxide, a first reaction chamber for reacting a metal silicate with acaustic material to produce a hydroxide of the metal, and a secondreaction chamber for contacting the metal hydroxide with the gas streamcontaining the carbon dioxide to produce a carbonate of the metal.

The invention also relates to a system for carbonating a metal silicate.The system comprises: (a) a supply of the metal silicate entering thesystem; (b) a source of carbon dioxide entering the system; (c) areactor structured for converting the metal silicate to a metalcarbonate and silica with the use of a caustic material, and With theuse of the carbon dioxide; and (d) the metal carbonate and the silicaexiting the system as separate products.

The invention also relates to a system for recovering a useful metalfrom rock. The system comprises: (a) a supply of rock entering thesystem, the rock containing the useful metal and a metal silicate; (b) asource of carbon dioxide entering the system; (c) a reactor structuredfor converting the metal silicate to a metal carbonate, with the use ofa caustic material, and with the use of the carbon dioxide; (d)apparatus for removing the useful metal from the rock; (e) a stream ofthe metal carbonate exiting the system; and (f) a stream of the usefulmetal exiting the system.

The invention also relates to a process of carbonating a metal silicate.The process comprises the steps of: (a) reacting the metal silicate witha caustic material to produce a hydroxide of the metal; (b) reacting themetal hydroxide with a source of carbon dioxide to produce a carbonateof the metal and to produce reconstituted caustic material; and (c)introducing the caustic material from step (b) into step (a).

The invention also relates to a process of carbonating a metal silicate.The process comprises reacting at least the metal silicate and a sourceof carbon dioxide to produce a carbonate of the metal, wherein thereaction is conducted at a pressure not greater than about 50 bars abovethe vapor pressure of pure water for the temperature of the reaction.

The invention also relates to a process of carbonating a metal silicate.The process comprises the steps of (a) reacting the metal silicate witha caustic alkali-metal hydroxide to produce a hydroxide of the metalformerly contained in the silicate; and (b) reacting the metal hydroxidewith a source of carbon dioxide to produce a carbonate of the metalformerly contained in the metal silicate of step (a).

The invention also relates to a process of producing a metal carbonate.The process comprises reacting an alkaline-earth metal hydroxide with atleast one of an alkali-metal carbonate, an alkali-metal bicarbonate, andcarbon dioxide, to produce a carbonate of the metal formerly containedin the metal hydroxide.

The following detailed discussion will make the advantages of theinvention apparent to the informed reader.

DETAILED DESCRIPTION OF THE INVENTION

Many different types of metal-silicate feedstocks are amenable tocarbonation by the invented process, including naturally occurringsilicates such as those present in rocks and clay-rich formations, aswell as silicates present in industrial waste products such as fly ashand waste concrete. Typically, the metal-silicate feedstock is composedof one or more calcium silicates, magnesium silicates, iron-bearingsilicates (such as basalt), or mixtures thereof, although other types ofsilicates can also be used. Some nonlimiting examples of these silicatesare described below. (Silicate feedstocks are referred to collectivelyherein as “metal silicates” with the understanding that this designationincludes any natural or man-made material, in the crystalline oramorphous state, that contains at least one metal along with silicon. Bythis definition, aluminosilicates are metal silicates because theycontain a metal, aluminum, along with silicon.)

Calcium silicates include wollastonite (CaSiO₃), calcic plagioclase[e.g., anorthite [(Ca_(x), Na_(1−x))(Al_(1+x)Si_(3−x)O₈, where x≧0.9,and labradorite [(Ca_(x), Na_(1−x))(Al₁₊, Si_(3−x)O₈, where 0.5≦x≦0.7],calcium-rich fly ash, basalt (a volcanic rock rich in Ca, Mg and Fe),calcium-rich montmorillonite [nominally (½Ca, Na)_(0.7)(Al, Mg,Fe)₄[(Si, Al)₈O₂₀] (OH)₄.nH₂O] and waste concrete. Contact-metamorphosedlimestones frequently contain wollastonite, large quantities of calcicplagioclase are present in many different types of common rocks, andcalcium-rich montmorillonite is found in special types of clay deposits.Basalt is a common rock in many terrestrial locations, and on the floorsof the world's major oceans.

Magnesium-rich silicates include olivine (specifically forsterite,Mg₂SiO₄), serpentine [Mg₃Si₂O₅(OH)₄], and basalt. Significant masses ofolivine- and serpentine-bearing rocks exist around the world,particularly in ultramafic complexes, and in large serpentinite bodies.

Iron-bearing silicates include fayalite, Fe₂SiO₄, and various naturalglasses (e.g., basaltic glass).

The metal silicates used as feedstocks for the process can have a widerange of initial particle sizes. Typically, it is desirable to reducethe particle size of the metal silicate(s) prior to chemical treatment.For example, the particle size of the metal silicate(s) may be reducedto an average diameter of less than about 100 microns. Any suitableequipment can be used to reduce particle size.

The process can optionally be conducted without heat pretreatment of themetal silicate feedstock.

In a preferred embodiment of the invention, one or more metal silicatesare transformed to one or more solid hydroxides by reaction with acaustic alkali-metal hydroxide, such as caustic soda (NaOH), in aqueoussolution. This is the first step of a preferred process; i.e., the metalsilicate(s) react with a caustic alkali-metal hydroxide to produce ahydroxide of the metal formerly contained in the silicate. This initialreaction is usually followed by physical and chemical segregation of theproduced solid(s) and “depleted” caustic liquid. In addition, it may bedesirable to separate the solid metal hydroxide(s) from any residualsolid silicate and/or oxide material that forms as a byproduct ofcaustic digestion.

Any suitable concentration of the caustic alkali-metal hydroxide inaqueous solution can be used to decompose the metal-silicate feedstock,including highly concentrated and very dilute solutions. The causticsolution is typically fairly concentrated, comprising, by weight, fromabout 30% to about 80% NaOH and from about 20% to about 70% water.

In the final step of the preferred process, the metal hydroxide formedin the first step is reacted with alkali-metal carbonate (e.g., sodiumcarbonate) and/or alkali-metal bicarbonate (e.g., sodium bicarbonate) toproduce a carbonate of the metal formerly contained in the metalsilicate. This reaction can be induced at any suitable set oftemperature-pressure conditions.

An intermediate step in the preferred process involves reacting carbondioxide with caustic alkali-metal hydroxide (e.g., NaOH) and/or alkalimetal silicate (e.g., Na₂SiO₃) to produce alkali-metal carbonate (e.g.,Na₂CO₃) and/or alkali-metal bicarbonate (e.g., NaHCO₃), ±water and/orsilica in either gelatinous or solid form. This step may or may not befollowed by precipitation of the Na₂CO₃ and/or NaHCO₃, which could beachieved by shifting the pH of the aqueous solution, or by evaporatingoff some of the water present.

When all of the foregoing steps are carried out using straight flue gasas a source of CO₂, capture, separation and sequestration of that gas isachieved in a single, integrated operation.

Advantageously, the caustic material produced in the intermediate stepof the process can usually be recycled back into the first step of theprocess. Thus, more generally, the invention relates to a process ofcarbonating a metal silicate which comprises the steps of: (a) reactingthe metal silicate with a caustic material to produce a hydroxide of themetal; (b) reacting the metal hydroxide with a source of carbon dioxideto produce a carbonate of the metal and to produce reconstituted causticmaterial; and (c) introducing the caustic material from step (b) intostep (a). The caustic material can be a caustic alkali-metal hydroxideor any other suitable caustic material.

It has been discovered by the inventors that the intermediate and finalsteps of the process can be conducted at a pressure not greater thanabout 50 bars above the vapor pressure of pure water for the temperatureof these two steps, typically not greater than about 30 bars, and moretypically not greater than about 20 bars, and often not greater thanabout 10 bars. The initial step can be conducted at a pressure slightlybelow the vapor pressure of pure water for the temperature of that step.Achieving rapid chemical reaction at low pressure is a key technologicaladvantage because relatively thin-walled pressure chambers will sufficeto safely contain the aqueous liquids (t gas) as reaction proceeds. Thiswill reduce the costs of commercial reactors built to implement theprocess on an industrial scale. Moreover, when total pressure is equalto the vapor pressure of the liquid phase, no investments in expensivepressure-intensifying equipment are required. On the other hand, higherfluid (liquid and/or gas) pressures at each step, particularly theintermediate step, could lead to more rapid and efficient chemicalreaction, in which case additional capital expenditures to make thecarbonation reactor more structurally robust, and to procure suitablepumping equipment, might be cost effective.

More generally, the invention relates to a process of carbonating ametal silicate. The process comprises reacting at least the metalsilicate and a source of carbon dioxide to produce a carbonate of themetal, wherein the reaction is conducted at a pressure not greater thanabout 50 bars above the vapor pressure of pure water for the temperatureof the reaction.

However, it may be beneficial to pressurize the CO₂-bearing gas to alevel above the vapor pressure of pure water for the temperature of thestep in which it is reacted, prior to, or during, production of metalcarbonate(s) and/or metal bicarbonate(s) in order to accelerate rates ofcarbonation. If CO₂ is captured, separated and liquified by a anotherprocess, then pressures up to ˜64 atm (the vapor pressure of pure liquidCO₂ at 25° C.) could be achieved simply by throttling flow of CO₂ intothe pressure chamber used to achieve carbonation.

In each step of the preferred process, the extent to which aqueousliquids are agitated or stirred, and control of the proportions ofphases as reaction proceeds, can be varied. In general, chemicalreaction is accelerated by vigorously agitating or rapidly stirring thereactants as processing proceeds, and by maintaining high fluid/solidratios.

More generally, the invention relates to a process of carbonating ametal silicate which comprises the steps of: (a) reacting the metalsilicate with a caustic alkali-metal hydroxide to produce a hydroxide ofthe metal formerly contained in the silicate; and (b) reacting the metalhydroxide with a source of carbon dioxide to produce a carbonate of themetal formerly contained in the metal silicate of step (a). Theinvention also relates to a process of producing a metal carbonate. Thisprocess comprises reacting an alkaline-earth metal hydroxide with atleast one of an alkali-metal carbonate, an alkali-metal bicarbonate, andcarbon dioxide, to produce a carbonate of the metal formerly containedin the metal hydroxide.

Some examples of reaction pathways for carbonating particular kinds ofmetal silicates are described below. It should be recognized that theinvention is not limited to these specific examples.

The invented process is advantageously used to carbonate calciumsilicates. For wollastonite (nominally CaSiO₃), the processing reactionsare:

(aq=aqueous, ↓=precipitate, liq=liquid), respectively. For anorthite(CaAl₂Si₂O₈, the processing reactions are:

(Other Ca-rich aluminosilicates carbonate by reactions similar to thosefor anorthite, with concomitant formation of various kinds of residualsolid materials.) The NaOH “regenerated” in the third step of theprocess (e.g., Reactions 7 and 10) can be recycled to decomposeadditional calcium silicate in step 1 (e.g., Reactions 5 and 8), or toform additional Na₂CO₃ in step 2 (e.g., Reactions 6 and 9). Table 1presents the results of experiments, performed by the inventors, thatconfirm Reactions 5-10.

In step 1 of the calcium-silicate carbonation processes (e.g., Reactions5 and 8), a solid, calcium-rich silicate feedstock is decomposed in anaqueous solution of caustic soda to produce crystalline sodium-calciumhydroxysilicate, or crystalline portlandite and crystallinehydroxy-sodalite. This step is followed by physical and chemicalsegregation of the precipitated solid(s) and “depleted” caustic liquid.In step 2 (e.g., Reactions 6 and 9), carbon dioxide is injected into aNaOH-bearing aqueous liquid, creating aqueous Na₂CO₃ and water. In step3 (e.g., Reactions 7 and 10), the Na₂CO₃ generated in step 2 is reactedwith the crystalline hydroxide(s) formed in step 1 to produce aqueousNaOH; and either crystalline calcite+crystalline pectolite, orcrystalline calcite+crystalline carbonate-cancrinite. These solidreaction products are thermodynamically stable, environmentally benign,and sparingly soluble in meteoric water.

The invented process also converts magnesium-rich silicates tomagnesium-rich carbonates. For example, the following process reactionsproduce magnesite (MgCO₃) from the magnesium-rich minerals olivine(specifically forsterite, Mg₂SiO₄) and serpentine [Mg₃Si₂O₅(OH)₄]:

[Net reaction: Mg₂SiO₄+2CO₂→2MgCO₃(↓)+SiO₂(→)], and

[Net reaction: Mg₃Si₂O₅(OH)₄+3CO₂→3MgCO₃(↓)+2SiO₂(↓)+2H₂O(liq)]. Table 2presents the results of experiments, performed by the writers, thatconfirm Reactions 11-17.

A key observation concerning Reaction 14 is that no heat pretreatment ofthe serpentine is required to achieve rapid and efficient production ofMg(OH)₂. This contrasts sharply with the so-called “direct” method forcarbonating serpentine (by the reactionMg₃Si₂O₅(OH)₄+3CO₂→3MgCO₃+2SiO₂+2H₂O), which requires heat pretreatmentof the serpentine at ˜600° C. to drive off structurally bound water.This extra step is necessary in the direct method of carbonatingserpentine because water-bearing (hydroxylated) serpentine reactssluggishly with CO₂, whereas dewatered (dehydroxylated) serpentine ishighly reactive. Dehydroxylation of serpentine makes the directcarbonation method very energy intensive and costly. In this regard, itis also noteworthy that the invented process, as applied to eitherolivine or serpentine, completely conserves the “rock solvent” (e.g.,NaOH), which lowers overall processing costs. By contrast, in the directmethod for carbonating serpentine and olivine, the rock solvent is(effectively) compressed, supercritical CO₂, which is expensive tocreate due to the high capital and operating costs of the mechanicalpumping that is required to achieve pressures as high as 185 atm.

Another important discovery made by the inventors is that Mg(OH)₂ can bereacted with Na₂CO₃ at 200° C. and elevated CO₂ fugacities to form thedouble carbonate eitelite [Na₂Mg(CO₃)₂] (see Table 2), which containstwice as much CO₂ as magnesite (MgCO₃). Crystallization of eitelitewould permit twice as much CO₂ to be sequestered per ton of minedMg-rich rock. In addition, like magnesite, eitelite is thermodynamicallystable, environmentally neutral, and only weakly soluble in meteoricwater. However, a disadvantage of eitelite crystallization is thatcaustic soda is not “regenerated” simultaneously, as it is whenmagnesite crystallizes (see Reactions 13 and 16). Therefore, when eitherolivine or serpentine is used to produce eitelite, the cost-savingsachieved by reducing the tonnage of mined Mg-rich rock are offset to asignificant extent by the continuous need to replenish the supply ofNaOH that is used to decompose the metal-silicate feedstock.

It should be clearly understood that the chemical formulae for thesolutes (substances dissolved in aqueous solution) in Reactions 11-13and 14-16 (specifically NaOH, Na₂SiO₃ and NaHCO₃) representstoichiometric components in aqueous solution, not “real” aqueousspecies. This convention was adopted (see also Reactions 5-7 and 8-10)for the sake of generality and simplicity. The particular species inaqueous solution created by our process (presently unknown) are ofconsiderable scientific interest; however, they need not be representedexplicitly in sets of process reactions such as those above, because thesolids that form and disappear in each process reaction, as well as thenet carbonation reaction for each metal silicate, do not depend on thechemical formulae that are used to represent the compositions ofsolutes. A simple example serves to illustrate this point. In Reactions11-13, the stoichiometric components NaOH, Na₂SiO₃ and NaHCO₃ can bereplaced by the ionic species OH⁻, SiO(OH)₃ ⁻, and HCO₃ ⁻, with sodiumion omitted because it is neither consumed nor produced in any reaction.This leads to the following alternative carbonation pathway forforsteritic olivine:

Mg₂SiO₄OH⁻+3H₂O→2Mg(OH)₂+SiO(OH)₃ ⁻  (18)

SiO(OH)₃ ⁻+CO₂→HCO₃ ⁻+SiO₂(↓)+H₂O  (19)

Mg(OH)₂+HCO₃ ⁻→MgCO₃(↓)+OH⁻+H₂O  (20)

Mg(OH)₂+CO₂→MgCO₃(↓)+H₂O  (21)

[Net reaction: Mg₂SiO₄+2CO₂→+2MgCO₃(↓)+SiO₂(↓)]. Comparing Reactions18-21 with Reactions 11-13, it is evident that the solids consumed andproduced, and the net reaction, are identical. Therefore, it should beclearly understood that the scope of our process for carbonating metalsilicates includes various self-consistent sets of reactions—i.e., setsof reactions involving the same solids, with metal silicate digestion byone or more caustic metal hydroxides (such as NaOH)—wherein solutes arerepresented by aqueous species of varying composition and charge, ratherthan by stoichiometric components.

The invented process may also convert iron-bearing silicates toiron-bearing carbonates, following carbonation pathways similar to thosedescribed above for Ca- and Mg-rich silicates. This is illustrated belowfor the iron-bearing silicate fayalite, Fe₂SiO₄:

[Net reaction: Fe₂SiO₄+2CO₂→2FeCO₃(↓)+SiO₂(↓)]. It is likely thatsimilar reactions will convert Fe-bearing silicate glasses (e.g.,basaltic glass) to one or more Fe-bearing carbonates.

In another embodiment of the invention, Ca-rich silicates are carbonatedin a single step. In this alternative carbonation pathway, the silicatefeedstock, aqueous NaOH, and Na₂CO₃ are reacted to produce crystallinecalcite±crystalline sodium-calcium hydroxysilicate±crystallinehydroxy-sodalite±crystalline carbonate-cancrinite±residual silicatematerial. For example, using a concentrated aqueous solution of causticsoda (such as 50 weight percent NaOH in H₂O) as the silicate solvent,and by adding abundant sodium carbonate to that solution to serve as asource of CO₂, the aluminum-bearing calcium silicates anorthite andlabradorite react with NaOH and Na₂CO₃ to form crystallinecalcite±crystalline hydroxy-sodalite±residual aluminosilicate materialat 200° C. and a total (fluid) pressure <15 atm (see Table 1).

The single-step process for carbonating Ca-rich silicates is simplerthan the three-step process described earlier, because physical andchemical segregation of the solids and liquids is not required. However,the one-step process has the disadvantage that carbonate-cancrinite istypically not among the solids that are produced. (In the “one-step”experiments performed by the inventors using Ca-rich silicates as asource of calcium, carbonate-cancrinite was only observed in experimentsperformed with Ca-rich fly ash.) A plausible explanation for this isthat the aqueous fluids produced in one-step experiments were typicallytoo basic to allow carbonate-cancrinite to crystallize. Therefore, in anindustrial-scale implementation of the one-step process, it may bedesirable, at some point, to lower the pH of the aqueous fluid toincrease the possibility that carbonate-cancrinite will be produced. Theformation of carbonate-cancrinite significantly increases the total“CO₂-loading” of the solids produced, and therefore leads to moreefficient and cost-effective carbonation of the Ca-silicate feedstock.

In still another embodiment, the invention provides a means forcarbonating magnesium and iron silicates in two steps. In step 1, themetal silicate(s) is (are) converted to Mg(OH)₂ and/or ironhydroxides(s)+Na₂SiO₃±SiO₂ by reaction with caustic soda in aqueoussolution (e.g., Reactions 11, 14 and 22). When this conversion isessentially complete, carbonation of Mg(OH)₂ and/or iron hydroxide(s) is(are) achieved by injecting CO₂ into the aqueous solution to form NaHCO₃(±Na₂CO₃)+silica gel and/or solid silica (step 2). MgCO₃ is formed whenthe Mg(OH)₂ produced in step 1 reacts with NaHCO₃ (±Na₂CO₃) and/oraqueous CO₂ (e.g., Reactions 13, 16 and 24).

Physical and/or chemical segregation of solids and liquids is notrequired in the two-step process for carbonating Mg- and Fe-richsilicates; therefore, it is intrinsically simplier than the three-stepprocess for carbonating forsteritic olivine (Reactions 11-13), and thefour-step process for carbonating serpentine (Reactions 14-17). On theother hand, the two-step process has the disadvantage that silica andmagnesite are produced simultaneously. It should be recognized thatReactions 11-13 and 14-17 produce separate “streams” of gelatinous/solidsilica and crystalline magnesite, which allows each substance to be usedas a feedstock for various commercial applications. In contrast, thetwo-step process for carbonating magnesium and iron silicates generatesa single stream of solid material, consisting of intimately mixedgelatinous/solid silica and crystalline magnesite. This creates twodifficulties. First, gelatinous silica (if formed) is readily mobilizedby meteoric water, and is therefore an inherently undesirable wasteproduct for near-surface terrestrial disposal. Second, when silica (inany form) and magnesite are intimately mixed, the high cost ofseparation effectively precludes any value-added commercial applicationsof each substance. Therefore, compared to separate masses of itsindividual components, an intimate mixture of silica and magnesite is aneconomic and environmental liability.

The invention also relates to a system for sequestering carbon dioxidefrom a gas stream. The system includes a gas stream containing carbondioxide, for example, a flue gas containing carbon dioxide or a streamof pure carbon dioxide gas. The system also includes a first reactionchamber for reacting a metal silicate with a caustic material to producea hydroxide of the metal. The caustic material can be a causticalkali-metal hydroxide, as discussed above, or any other causticmaterial suitable for the reaction. The system also includes a secondreaction chamber for contacting the metal hydroxide with the gas streamcontaining the carbon dioxide to produce a carbonate of the metal. Thefirst and second reaction chambers can be located in the same reactor orin different reactors. More than two reaction chambers can optionally beused in the system.

The invention also relates to a system for carbonating a metal silicate.The system includes a supply of the metal silicate entering the system,and a source of carbon dioxide entering the system. The source of carbondioxide can be a gas stream containing carbon dioxide, or it can be acompound containing carbon dioxide, such as an alkali-metal carbonate orbicarbonate. The system also includes a reactor structured forconverting the metal silicate to a metal carbonate and silica with theuse of a caustic material, and with the use of the carbon dioxide. Anysuitable caustic material can be used, such as a caustic alkali-metalhydroxide. The reaction can include other reactants in addition to themetal silicate, the caustic material and the carbon dioxide. The metalcarbonate and the silica exit the system as separate products. In apreferred embodiment, the metal silicate is magnesium silicate and themetal carbonate is magnesite. Typically, the magnesite and the silicaexiting the system have a purity of at least about 90%. This iseconomically advantageous, and it contrasts with previously knownprocesses that produce mixtures of magnesite and silica.

The invention also relates to a system for recovering a useful metalfrom rock. The system includes a supply of rock entering the system, therock containing the useful metal and a metal silicate, and a source ofcarbon dioxide entering the system. The system also includes a reactorstructured for converting the metal silicate to a metal carbonate, withthe use of a caustic material, and with the use of the carbon dioxide.As discussed above, any suitable caustic material can be used, and thesource of carbon dioxide can be a gas containing carbon dioxide or acompound containing carbon dioxide. The system also includes apparatusfor removing the useful metal from the rock. The apparatus can belocated at any suitable location in the system. For example, magneticapparatus can be used for removing magnetite from serpentine prior tothe reactor. Alternatively, the reactor can produce the metal carbonateand a remaining rock portion, and the system can include apparatus forremoving the useful metal from the remaining rock portion at a locationsubsequent to the reactor. The system also includes a stream of themetal carbonate exiting the system, and a stream of the useful metalexiting the system. The streams can be in any form, such as truckloads,trainloads, or other means of conveying the metal carbonate and theuseful metal from the system.

The carbonation pathways for Ca-, Mg- and Fe-rich silicates describedabove are similar to, but distinctly different from, the commercialchemical processes that are used to extract alumina (Al₂O₃) from bauxite(aluminum ore), and to generate caustic soda from trona (a rock rich insodium carbonate, Na₂CO₃). In the treatment of bauxite ore by thewell-known Bayer Process, caustic soda is used to remove reactive silicaand iron oxides, and to dissolve aluminum oxides (gibbsite, boehmite anddiaspore). Dissolution of silica by the caustic solution produces sodiumsilicate (nominally Na₂SiO₃), which quickly reacts with sodium aluminate(NaAlO₂) to form crystalline hydroxy-sodalite [Na₈(AlSiO₄)₆(OH)₂]. Thisdesilication of the solution is detrimental to the overall processbecause it consumes caustic soda, and the total mass of dissolved sodiumaluminate is lowered. (With less sodium aluminate in solution, lesshigh-purity gibbsite is precipitated in a later stage of the process,and as a second consequence, less caustic soda is regenerated bygibbsite precipitation. Regenerated caustic soda is recycled in theBayer Process to treat additional batches of bauxite ore.)

While the Bayer Process has several characteristics in common with theprocess described in this document—in particular, the use of causticsoda to decompose the metal-bearing, solid feedstock, and regenerationof caustic soda at a subsequent stage of the process—two majordifferences are also evident. The first is that the goals of the twoprocesses are totally different: the Bayer Process was developed toproduce a solid concentrate rich in aluminum (precipitated gibbsite),whereas the principal intent of the invented metal-silicate carbonationprocess is to form stable metal-carbonate compounds for long-term CO₂sequestration. Due to this key difference, the “ore” used for the twoprocesses is much different. In the Bayer Process, deeply weathered,unconsolidated rock material is reacted because it is rich in aluminaand poor in silica. In the present invention, the silica content of themetal-silicate feedstock is not a significant factor, except that agreater silica content generally means a lower Ca, Mg and/or Fe content,which is undesirable. The second key difference is that, in the presentinvention, caustic soda is regenerated during crystallization of one ormore solid metal carbonates, whereas in the Bayer Process, caustic sodais regenerated during the production of high-purity gibbsite.

In the commercial process of reacting trona with lime (CaO) to producecaustic soda plus calcite, the intent is solely to produce caustic soda;the entire amount of calcite formed as a byproduct is subsequentlycalcined to regenerate lime, which is recycled for reaction withadditional batches of trona ore to produce more caustic soda. Thus,there is no CO₂ sequestration achieved in the commercial treatment oftrona. Moreover, trona contains little or no silica and alumina;consequently, its treatment to generate caustic soda does not consume orproduce any significant amount of silicate material.

The invented metal-silicate carbonation process may have the followingpractical uses and benefits:

CO₂ sequestration. The process was designed mainly to producecrystalline carbonates that persist indefinitely in most continentalsettings. The results of autoclave experiments presented in Tables 1 and2 demonstrate that, under suitable conditions: Ca-rich silicates arereadily converted to calcite (CaCO₃)±carbonate-cancrinite[Na₈(AlSiO₄)₆CO₃.2H₂O], and Mg-rich silicates are quickly transformed tomagnesite (MgCO₃) and/or eitelite [Na₂Mg(CO₃)₂]. These four crystallinecarbonate compounds bind CO₂ indefinitely in most non-acidic terrestrialenvironments, and are completely harmless to all flora and fauna.

Neutralization of highly acidic soils. Locally, calcium and magnesiumcarbonates can have commercial value as soil amendments. While addingcarbonates to highly acid soils to increase pH ultimately releases CO₂to the atmosphere, significant environmental and economic benefits wouldaccrue if the treated land was made more biologically productive. Theamended soils might be used to grow crops or trees, either of whichcould have a total carbon sequestration potential higher than that ofthe crystalline carbonate amendment. Using calcium and magnesiumcarbonates for this purpose would also lessen demand for lime producedby calcining limestone, and this would help lower CO₂ emissions to theatmosphere.

Recovery of useful metals. Many rock formations contain useful metals(e.g., iron, copper, nickel and platinum) at concentrations that arecurrently uneconomical to mine. If, however, mining and grinding werealready being performed to create a metal-silicate feedstock forcarbonation, one or more metals that are not carbonated could beextracted as a byproduct(s), thereby reducing the costs ofmetal-silicate carbonation.

Elimination of hazardous mine tailings. Mine tailings, consisting ofcrushed rock material from which metals or other valuable materials havebeen extracted—along with the “overburden” (soil and regolith) that isremoved to access buried ore horizons—are an important waste problem formany active and abandoned mines. Thus, it is significant that a nearbymineral carbonation reactor might be able to use them as a source ofcalcium and magnesium. This would enhance environmental restoration, andreduce the costs of mining metal silicate(s) for carbonation.

Production of high-purity silica. When olivine and/or serpentine is(are) carbonated by the invented process, the solid effluent produced instep 2 (Reactions 12 and 15) is high-purity silica, which can be refinedfor use in manufacturing silica-based desiccants, silica brick, siliconcarbide, and various types of glass. High-purity silica is also apotential source of elemental silicon the foundation material fornumerous semiconducting electronic devices. Finally, it may be possibleto use amorphous silica to form melanophlogite, a silica-rich compoundwith a cage structure that can accommodate as many as six CO₂ “guest”molecules for every 46 molecules of SiO₂. Creation of substantialamounts of melanophlogite would significantly increase the totalCO₂-loading of the solids generated by the invented process.

Production of high-purity magnesite. When olivine and/or serpentine is(are) carbonated by the invented process, the solid effluent produced instep 3 for olivine carbonation (Reaction 13) and in steps 3 and 4 forserpentine carbonation (Reactions 16 and 17) is high-purity magnesite,which can be used to produce magnesite cement.

Capture and separation of CO₂ from flue gas. Another potentialapplication of the invented process is especially important, as it wouldgreatly reduce costs associated with capture and separation of CO₂ atfossil fuel-fired power plants, cement factories, and steel mills.Specifically, step 2 in the process (e.g., Reactions 6, 9, 12 and 15)permits CO₂ to be captured and separated from flue gas by bubbling thegas through a NaOH- and/or Na₂SiO₃-bearing aqueous liquid. The CO₂ wouldbe transformed to aqueous±crystalline Na₂CO₃ and/or NaHCO₃, and thenitrogen-rich gas effluent could either be refined to producehigh-purity nitrogen for commercial use, or simply released harmlesslyto the atmosphere.

TABLE 1 Results of Experiments with Solid Calcium Silicates Date/Temperature/ XRD Solid Run Products, and Duration^(@) StartingMaterials^(@,#) Pressure Number Initial/Final Solution pH^(&) 7-10-02/anorthite + NaOH + H₂O 200° C./ LMAS001-2 portlandite +hydroxy-sodalite + 72 hrs <15 atm LMAS003 hydrogrossularite(?) 8-1-02/anorthite + NaOH + Na₂CO₃ + 200° C./ LMA059 calcite + hydroxy-sodalite72 hrs H₂O <15 atm 10-4-02/ prereacted anorthite + 200° C./ LMA070calcite + carbonate-cancrinite, 72 hrs Na₂CO₃ + H₂O <15 atm pH: 12.5/≧148-20-02/ labradorite + NaOH + H₂O 200° C./ LMA060 portlandite +hydroxy-sodalite + 72 hrs <15 atm cancrinite + hydrogrossular(?)8-28-02/ labradorite + NaOH + 200° C./ LMA061 calcite +hydroxy-sodalite + 72 hrs Na₂CO₃ + H₂O <15 atm hydrogrossular + minorAUCP, pH: 12.5/≧14 10-8-02/ prereacted labradorite + 200° C./ LMA071calcite + carbonate-cancrinite, 72 hrs Na₂CO₃ + H₂O <15 atm pH: 12.5/≧1411-8-02/ Ca(OH)₂ + Mg(OH)₂ + 200° C./ LMA077 calcite + brucite, pH:12.5/≧14 72 hrs Na₂CO₃ + H₂O <15 atm 11-26-02/ wollastonite + NaOH + H₂O200° C./ LMA081 sodium-calcium 72 hrs <15 atm hydroxysilicate 12-2-02/wollastonite + NaOH + 200° C./ LMA083 calcite + sodium-calcium 72 hrsNa₂CO₃ + H₂O <15 atm hydroxysilicate, pH: 12.5/≧14 12-10-02/ Ca-rich flyash + NaOH + 200° C./ LMA084 sodalite + AUCP 72 hrs H₂O <15 atm12-13-02/ Ca-rich fly ash + NaOH + 200° C./ LMA085 calcite +carbonate-cancrinite + 72 hrs Na₂CO₃ + H₂O <15 atm sodalite + AUCP, pH:12.5/≧14 12-17-02/ Ca(OH)₂ + Na₂CO₃ + H₂O 200° C./ LMA089 calcite, pH:12.5/≧14 72 hrs <15 atm 1-3-03/ prereacted wollastonite + 200° C./LMA094 calcite + pectolite, pH: 72 hrs NaOH + Na₂CO₃ + H₂O <15 atm12.5/≧14 9-16-02/ basalt + NaOH + H₂O 200° C./ LMA064 portlandite +sodalite + brucite 72 hrs <15 atm 9-30-02/ basalt + NaOH + Na₂CO₃ + 200°C./ LMA063 brucite + AUCP 72 hrs H₂O <15 atm 10-14-02/ prereactedbasalt + NaOH + 200° C./ LMA059 calcite + carbonate-cancrinite, 72 hrsNa₂CO₃ + H₂O <15 atm pH: 12.5/≧14 ^(@)Sources of solid-silicate startingmaterials: wollastonite - Willsboro, NY; anorthite - Grass Valley, CA;labradorite - Nain, Labrador, Canada; basalt (USGS standard BCR-2) -Portland, OR; Ca-rich fly ash, Joseph City, AZ; Ca(OH)₂ (portlandite) -commercially manufactured, reagent-grade chemical compound; Mg(OH)₂(brucite) - commercially manufactured, reagent-grade chemical compound.NaOH (caustic soda) was added to each unreacted starting sample as a 50weight percent solution of NaOH in H₂O. ^(#)For each solid silicate,“prereacted” means reaction with a 50 weight percent solution of NaOH inH₂O at 200° C., P < 15 atm, before the experiment was performed. In allcases, this step had the effect of converting the silicates tocrystalline hydroxide(s) prior to reaction with Na₂CO₃ to form one ormore crystalline carbonate compounds. ^(&)Chemical compositions ofcrystalline reaction products: calcite—CaCO₃; portlandite—Ca(OH₂);hydroxy-sodalite—Na₈(AlSiO₄)₆(OH)₂; sodium-calciumhydroxysilicate—NaCaSiO₃(OH); carbonate-cancrinite—Na₈(AlSiO₄)₆CO₃•2H₂O;brucite—Mg(OH)₂; pectolite—NaCa₂Si₃O₈(OH);hydrogrossular—Ca₃Al₂(SiO₄)₃•xH₂O. AUCP = additional unidentifiedcrystalline phase(s).

TABLE 2 Results of Experiments with Crystalline Magnesium SilicatesDate/ Temperature/ XRD Solid Run Products, and Duration^(@) StartingMaterials^(@,#) Pressure Number Initial/Final Solution pH^(&) 6-10-02/serpentine + NaOH + H₂O 200° C./ LMAS001-1 brucite + relict serpentine72 hrs <15 atm 7-1-02/ olivine + NaOH + H₂O 200° C./ — brucite 72 hrs<15 atm 7-16-02/ prereacted olivine + Na₂CO₃ + 200° C./ LMAS006 brucite72 hrs H₂O <15 atm 9-20-02/ serpentine + NaOH + H₂O 200° C./ LMA069brucite + trace AUCP 24 hrs <15 atm 9-23-02/ serpentine + NaOH + H₂O200° C./ LMA067 brucite + trace AUCP 72 hrs <15 atm 9-25-02/ Mg(OH)₂ +Na₂CO₃ + H₂O 200° C./ LMA072 brucite, pH: 12.5/12.5 72 hrs <15 atm10-17-02/ serpentine + NaOH + H₂O 22° C./ LMA073 serpentine + brucite 72hrs 1 atm 10-18-02/ serpentine + NaOH + H₂O 22° C./ LMA075 brucite +minor relict 72 hrs at 1 atm, serpentine 22° C., then 3 hrs 200° C./ at200° C. <15 atm 11-15-02/ Mg(OH)₂ + Na₂CO₃ + H₂O 300° C./ LMA078 brucite72 hrs <86 atm 11-19-02/ Mg(OH)₂ + Na₂CO₃ + H₂O 375° C./ LMA079 brucite~1 hr <220 atm 12-23-02/ Mg(OH)₂ + Na₂CO₃ + H₂O 200° C./ LMA088eitelite + minor relict brucite 72 hrs ~60 atm* 12-30-02/ Mg(OH)₂ + H₂O200° C./ LMA090 magnesite + minor relict 72 hrs ~60 atm* brucite, finalpH: 5.5 1-17-03/ Mg(OH)₂ + Na₂CO₃ + H₂O 300° C./ LMAS014 brucite +eitelite, pH: 11.5-12/ 72 hrs ~60 atm* 12 3-17-03/ Mg(OH)₂ + Na₂CO₃ +H₂O 100° C./ LMAS017 brucite + magnesium carbonate 72 hrs <1 atm hydrate3-21-03/ Mg(OH)₂ + Na₂CO₃ + H₂O 80° C./ LMAS022 brucite + magnesiumcarbonate 72 hrs <1 atm hydrate, pH: 12.5/12.5 4-07-03/ Mg(OH)₂ +Na₂CO₃ + H₂O 125° C./ LMAS030 brucite + magnesium carbonate 72 hrs <2atm hydrate, final pH: 11.5 4-16-03/ Mg(OH)₂ + NaHCO₃ + H₂O 200° C./ —magnesite + minor eitelite + 72 hrs <15 atm minor relict brucite^(@)Sources of crystalline silicate starting materials: olivine - TwinSisters Peak, WA; serpentine (antigorite variety) - Cedar Hill Quarry,Lancaster County, PA; Mg(OH)₂ (brucite) - commercially manufactured,reagent-grade chemical compound. NaOH (caustic soda) was added to eachunreacted starting sample as a 50 weight percent solution of NaOH inH₂O. ^(#)For each solid silicate, “prereacted” means reaction with a 50weight percent solution of NaOH in H₂O at 200° C., P < 15 atm, beforethe experiment was performed. In all cases, this step had the effect ofconverting olivine or serpentine to brucite prior to reaction withNa₂CO₃ or NaHCO₃ to form one or more crystalline carbonate compounds. *Atotal (fluid) pressure of approximately 60 atm was achieved by injectingCO₂ into the headspace of the autoclave, and keeping that spaceconnected to an external cylinder filled with liquid CO₂ at roomtemperature. ^(&)Chemical compositions of crystalline reaction products:olivine—Mg₂SiO₄; serpentine—Mg₃Si₂O₅(OH)₄; brucite—Mg(OH)₂;eitelite—Na₂Mg(CO₃)₂. magnesium carbonate hydrate—MgCO₃•xH₂O. AUCP =additional unidentified crystalline phase(s).

In accordance with the provisions of the patent statutes, the principleand mode of operation of this invention have been explained andillustrated in its preferred embodiments. However, it must be understoodthat this invention may be practiced otherwise than as specificallyexplained without departing from its spirit or scope.

1. A process of sequestering carbon dioxide comprising the steps of: (a)reacting a metal silicate with a first caustic alkali-metal hydroxide toproduce a metal hydroxide comprising metal formerly contained in themetal silicate; (b) reacting carbon dioxide with at least one of asecond caustic alkali-metal hydroxide and an alkali-metal silicate toproduce at least one of an alkali-metal carbonate and an alkali-metalbicarbonate; and (c) reacting the metal hydroxide product of step (a)with at least one of the alkali-metal carbonate and the alkali-metalbicarbonate produced in step (b) to produce a metal carbonate comprisingmetal formerly contained in the metal silicate of step (a).
 2. A processaccording to claim 1, wherein the reaction of step (c) further producesa reconstituted caustic alkali-metal hydroxide, and wherein the processfurther comprises an additional step (d) of recycling the reconstitutedcaustic alkali-metal hydroxide from step (c) into the reaction of step(a).
 3. A process according to claim 1, wherein steps (b) and (c) areconducted at a pressure not greater than about 50 bars above the vaporpressure of pure water.
 4. A process according to claim 1, wherein whenthe metal silicate is magnesium silicate the process produces at leastone of magnesite and eitelite, or when the metal silicate is a calciumsilicate the process produces calcite, or when the metal silicate is aniron-bearing silicate the process produces siderite.
 5. A processaccording to claim 1, wherein the metal silicate comprises calciumsilcates, magnesium silicates, iron-bearing silicates, or mixturesthereof, and at least one of the first caustic alkali-metal hydroxideand the second caustic alkali-metal hydroxide is comprised of sodiumhydroxide, potassium hydroxide, lithium hydroxide, or a mixture thereof.6. A process according to claim 1, further comprising reducing theparticle size of the metal silicate to an average diameter of less thanabout two millimeters.
 7. A process of producing a metal carbonatecomprising: reacting an alkaline-earth metal hydroxide with at least oneof an alkali-metal carbonate, an alkali-metal bicarbonate, and carbondioxide, to produce a metal carbonate comprising metal formerlycontained in the metal hydroxide.
 8. A process of carbonating a metalsilicate comprising: reacting at least the metal silicate and a sourceof carbon dioxide to produce a carbonate of the metal, wherein thereaction is conducted at a pressure not greater than about 50 bars abovethe vapor pressure of pure water for temperature of reaction.
 9. Aprocess according to claim 8, wherein the reaction is conducted at apressure not greater than about 30 bars above the vapor pressure of purewater for the temperature of the reaction.